Intermolecular Forces, Liquids

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Chapter 11
Intermolecular Forces, Liquids Solids
2
Overview

• Liquids Solids
• Intermolecular Forces
• Liquids
• Phase Changes
• Vapor Pressure
• Phase Diagrams
• Solids -- Structure
• Bonding Types in Solids

3
Liquids Solids

• Solids
• particles close together
• locked into relative positions (crystalline)
• strong interactions (interparticle forces)
• Liquids
• particles farther apart
• mobile relative to each other
• weaker interactions between particles

Remember
4
D Hcondensation -
D Hvaporization
condensation
- E
E
vaporization
Gas
Liquid
5
D Hfreezing -
D Hmelting
freezing
- E
E
melting
Liquid
Solid
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Intermolecular Forces

• Strength of IM Forces determine boiling points
  and melting points
• Ion-Dipole Forces
• occur between ions and dipoles
• between charged particles and neutral, polar
  covalent particles
• Dipole-Dipole Forces
• occur between two dipoles
• between two, neutral, polar covalent particles

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Ion - Dipole Force (40 - 600 kJ/mol)
d
H

d -
d
O
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Dipole - Dipole Force (5 - 25 kJ/mol)
d
d
d -
d -
Hydrogen Bonding Force (4 - 25 kJ/mol)
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• Hydrogen Bonding Forces
• special case of Dipole-Dipole forces
• occur between two, neutral, polar covalent
  particles which
• have a H atom (which is bound to an O, F or N
  atom) for the () dipole
• have an O, F, or N atom for the (-) dipole
• extra strength due to
• small size and large EN of O, F or N
• and small size of H

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• London Dispersion Forces
• occur between all particles even neutral,
  non-polar covalent particles
• occur between an instantaneous dipole and an
  induced dipole
• force is weak but strengthens with increasing
  polarizability of the particles
• polarizability of the particles increases with
  increasing size or mass

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Dipole - Induced Dipole (2 - 10 kJ/mol)
d
d -
d
d -
polar molecule non-polar molecule
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Induced Dipole - Induced Dipole
Force (0.05 - 40 kJ/mol)
d -
d
d -
d
London Dispersion
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Strength of Forces

• Ion - Dipole
• Hydrogen Bonding
• Dipole - Dipole
• Dipole - Induced Dipole
• London Dispersion

Increasing
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Boiling Points of Noble Gases
Xe 166.1
Kr 120.9
Ar 87.5
Temp K
Ne 27.3
He 4.6
MM
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Boiling Points of Halogen Diatomics
I2 457.6
Temp K
Br2 332.0
Cl2 238.6
F2 85.1
MM
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Boiling Points of Group 6 Dihydrides
H2O 373
H2Te 271
Temp K
H2Se 231
SnH4
H2S 212
GeH4 184
SiH4 161
CH4 109
MM
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What Does it Mean?

• hydration is very important in the solvation
  process compound formation
• stronger the interaction, the more energy is
  released, more exothermic
• the strength of the interaction determines the
  state of the substance
• unusual properties of water are due to
  hydrogen bonding

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What kind of forces are between

• O2 molecules
• H2O and NH3 molecules
• Ne atoms
• HF and NH3 molecules
• CH4 and Br2 molecules

London Dispersion Hydrogen Bonding London
Dispersion Hydrogen Bonding London Dispersion
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Properties of Liquids

• Viscosity
• resistance to flow
• ease with which liquid particles move relative to
  one another
• related to attractive forces between particles
• and structural properties of the particles
  themselves
• decreases with increasing energy (temperature) of
  particles

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• Surface Tension
• energy required to increase the surface area of a
  liquid by a unit amount
• a sphere produces the minimum surface area
• competition between cohesive forces vs adhesive
  forces
• cohesive forces tend to minimize surface area
• adhesive forces tend to maximize surface area
• high surface tension reflects strong cohesive
  forces
• capillary action -- low surface tension, strong
  adhesive forces

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Phase Changes

• Energy Changes
• melting, solid to liquid
• endothermic -- D Hfusion
• vaporization, liquid to gas
• endothermic -- D Hvaporization
• condensation, gas to liquid
• exothermic -- D Hcondensation - D
  Hvaporization
• freezing, liquid to solid
• exothermic -- D Hfreezing - D Hfusion

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Phase Changes, contd

• sublimation, solid to gas
• endothermic -- D Hsublimation
• deposition, gas to solid
• exothermic -- D Hdeposition - D Hsublimation

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Properties of Liquids
vaporization -- endothermic
liquid
gas
condensation -- exothermic
freezing -- exothermic
solid
liquid
melting -- endothermic
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constant temp
liquid gas in equilibrium
gas
Temp
solid liquid in equilibrium
liquid
solid
Time
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• Critical Temperature Pressure
• critical temp.
• highest temperature at which a substance can
  exist as a liquid
• critical pressure
• pressure required to cause liquefaction at the
  critical temperature
• a gas cannot be liquefied above the critical
  temperature
• critical point
• corresponds to Tc and Pc
• the point above which a supercritical gas exists
• the substance cannot be liquefied by increasing
  the pressure

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217.7 atm Pc
Critical Point
Liquid H2O
Vapor
Pressure
Tc 647.6 K
Temperature
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Vapor Pressure

• Molecular Description
• liquids have a distribution of energies for the
  liquid molecules
• at any temperature, some molecules have
  sufficient energy for vaporization
• the higher the temperature, the greater number of
  molecules with energy of vaporization
• at constant temperature, average energy of
  molecules is constant but in dynamic equilibrium
  
• vapor pressure is the pressure exerted by
  vaporized molecules when liquid and vapor states
  are in dynamic equilibrium

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gas
liquid
Equilibrium Vapor Pressure
Evaporation
Eq. Vapor Pressure -- partial pressure of gas
over a liquid at equilibrium
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• Volatility, Vapor Pressure Temp.
• substances with high vapor pressure are volatile
  
• in a open container, dynamic equilibrium cannot
  be established -- complete evaporation
• higher the temperature, greater vapor pressure,
  greater volatility
• Vapor Pressure Boiling Point
• boiling point -- temperature at which the vapor
  pressure atmospheric pressure
• boiling point increases with increasing external
  pressure, vice versa
• normal boiling point -- temp. at which the vapor
  pressure 1 atm

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Boiling Point temperature at which the vapor
pressure is equal to the external or atmospheric
pressure
Normal Boiling Point temperature at which the
vapor pressure is equal to 1.00 atmosphere
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Which has highest boiling point
H2O BrCl HCl C2H6

• H2O or H2S
• BrCl or Cl2
• BrCl or HCl
• CH4 or C2H6

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Calculate the Energy of Vaporization

• molar enthalpy of vaporization(kJ/mol) x
  amount of substance(moles)
• DHvap water 40.7 kJ/mol _at_ 100º
• Clausius Clapeyron Equation ln P -
  DHvap / RT C

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vapor
solid
Phase Diagrams
solid
liquid
liquid
vapor
normal freezing pt.
normal boiling pt.
liquid
vapor
Pressure
solid
H2O
Triple Point
Temperature
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Structures of Solids

• Unit Cells
• characteristic 3-dimensional repeating unit of a
  crystalline substance
• primitive cubic (sc) 1 atom/u.c.
• body-centered cubic (bcc) 2 atoms/u.c.
• face-centered cubic (fcc) 4 atoms/u.c.

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primitive cubic
body-centered cubic
face-centered cubic
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Close Packing of Spheres

• Layered as ABABABAB or
• ABCABC

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Physical Properties of Solids
melting -- endothermic
solid
liquid
freezing -- exothermic
sublimation -- endothermic
solid
gas
deposition -- exothermic
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Melting Point Temperature at which the crystal
lattice breaks down DHfusion - DHfreezing
melting
freezing
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Which one of each pair will have the higher
melting point?
NaF HCl Br2 NaCl H2O

• NaF or NaI
• HCl or HI
• O2 or Br2
• NaCl or BrCl
• H2S or H2O

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Types of Solids Types Units

• Ionic cations anions
• Metallic metal atoms
• Molecular molecules
• Network atoms
• Amorphous irregular networks

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Network Solid
Diamond
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Network Solid
Graphite
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Na
NaCl
Cl